View source for Galvanic cell

{{Short description|Electrochemical device}}[[File:Galvanic cell with no cation flow.svg|upright=1.4|thumb|Galvanic cell with no cation flow]]
A '''galvanic cell''' or '''voltaic cell''', named after the scientists [[Luigi Galvani]] and [[Alessandro Volta]], respectively, is an [[electrochemical cell]] in which an electric current is generated from [[Spontaneous reaction|spontaneous]] [[oxidation–reduction]] reactions. An example of a galvanic cell consists of two different metals, each immersed in separate beakers containing their respective metal ions in solution that are connected by a [[salt bridge]] or separated by a porous membrane.<ref name=":0">{{Cite book|last1=McMurry|first1=John|last2=Fay|first2=Robert C.|last3=Robinson|first3=Jill K.|title=Chemistry|publisher=Pearson|year=2015|isbn=978-0-321-94317-0|edition=7th|location=Boston|pages=762|oclc=889577526}}</ref>

Volta was the inventor of the [[voltaic pile]], the first [[Battery (electricity)|electrical battery]]. Common usage of the word ''battery'' has evolved to include a single galvanic cell, but the first batteries had many galvanic cells.<ref>[http://www.merriam-webster.com/dictionary/battery "battery" (def. 4b)], ''Merriam-Webster Online Dictionary'' (2008). Retrieved 6 August 2008.</ref>

==History==
In 1780, [[Luigi Galvani]] discovered that when two different metals (e.g., copper and zinc) are in contact and then both are touched at the same time to two different parts of a muscle of a frog leg, to close the circuit, the frog's leg contracts.<ref>{{ cite book | title = Daniell Cell | last = Keithley | first = Joseph F | publisher = John Wiley and Sons | year = 1999 | isbn = 978-0-7803-1193-0 | pages = 49–51 }}</ref> He called this "[[bioelectricity|animal electricity]]".  The frog's leg, as well as being a detector of electrical current, was also the [[electrolyte]] (to use the language of modern chemistry).

A year after Galvani published his work (1790), [[Alessandro Volta]] showed that the frog was not necessary, using instead a force-based detector and brine-soaked paper (as electrolyte). Earlier Volta had established the law of capacitance {{nowrap|1=''C'' = {{sfrac|''Q''|''V''}}}} with force-based detectors.  In 1799 Volta invented the voltaic pile, which is a stack of galvanic cells each consisting of a metal disk, an electrolyte layer, and a disk of a different metal. He built it entirely out of non-biological material to challenge Galvani's (and the later experimenter [[Leopoldo Nobili]])'s animal electricity theory in favor of his own metal-metal contact electricity theory.<ref>Kipnis, Nahum (2003) [http://ppp.unipv.it/Collana/Pages/Libri/Saggi/Nuova%20Voltiana5_PDF/p__143-162.pdf "Changing a theory: the case of Volta's contact electricity"], ''Nuova Voltiana'', Vol. 5. Università degli studi di Pavia, 2003 {{ISBN|88-203-3273-6}}. pp. 144–146</ref> [[Carlo Matteucci]] in his turn constructed a [[Frog battery|battery entirely out of biological material]] in answer to Volta.<ref>Clarke, Edwin; Jacyna, L. S. (1992) [https://books.google.com/books?id=38Sjkp-JlPcC&pg=PA198 ''Nineteenth-Century Origins of Neuroscientific Concepts''], University of California Press. {{ISBN|0-520-07879-9}}. p. 199</ref>  Volta's contact electricity view characterized each electrode with a number that we would now call the work function of the electrode.  This view ignored the chemical reactions at the electrode-electrolyte interfaces, which include {{chem2|H2}} formation on the more noble metal in Volta's pile.

Although Volta did not understand the operation of the battery or the galvanic cell, these discoveries paved the way for electrical batteries; Volta's cell was named an [[List of IEEE milestones|IEEE Milestone]] in 1999.<ref>{{cite web |url=http://www.ieeeghn.org/wiki/index.php/Milestones:Volta%27s_Electrical_Battery_Invention,_1799 |title=Milestones:Volta's Electrical Battery Invention, 1799 |work=IEEE Global History Network |publisher=IEEE |access-date=26 July 2011}}</ref>

Some forty years later, Faraday (see [[Faraday's laws of electrolysis]]) showed that the galvanic cell—now often called a voltaic cell—was chemical in nature.  Faraday introduced new terminology to the language of chemistry: [[electrode]] ([[cathode]] and [[anode]]), [[electrolyte]], and [[ion]] ([[cation]] and [[anion]]).  Thus Galvani incorrectly thought the source of electricity (or source of [[electromotive force]] (emf), or seat of emf) was in the animal, Volta incorrectly thought it was in the physical properties of the isolated electrodes, but Faraday correctly identified the source of emf as the chemical reactions at the two electrode-electrolyte interfaces. The authoritative work on the intellectual history of the voltaic cell remains that by Ostwald.<ref>{{ cite book |last=Ostwald |first=Wilhelm |date=1980 |title=Electrochemistry: History and Theory }}</ref><!--publisher = translated from German, New Delhi: Amerind Publishing Co.; Springfield, Va, 1980, 2 vols. year = 1896 Veit & Co., Leipzig-->

It was suggested by [[Wilhelm König]] in 1940 that the object known as the [[Baghdad battery]] might represent galvanic cell technology from ancient [[Parthia]].  Replicas filled with citric acid or grape juice have been shown to produce a voltage.  However, it is far from certain that this was its purpose—other scholars have pointed out that it is very similar to vessels known to have been used for storing parchment scrolls.<ref>Haughton, Brian (2007) ''Hidden History: Lost Civilizations, Secret Knowledge, and Ancient Mysteries''. Career Press. {{ISBN|1564148971}}. pp. 129–132</ref>

==Principles==
[[file:Galvanic cell labeled.svg|upright=1.55|thumb|Schematic of Zn–Cu galvanic cell]]

Galvanic cells are extensions of spontaneous [[redox]] reactions, but have been merely designed to harness the energy produced from said reaction.<ref name=":0"/> For example, when one immerses a strip of zinc metal (Zn) in an aqueous solution of copper sulfate ({{chem2|CuSO4}}), dark-colored solid deposits will collect on the surface of the zinc metal and the blue color characteristic of the {{chem2|Cu(2+)}} ion disappears from the solution. The depositions on the surface of the zinc metal consist of copper metal, and the solution now contains zinc ions. This reaction is represented by:
:{{chem2|Zn(s) + Cu(2+)(aq) -> Zn(2+)(aq) + Cu(s)}}

In this [[redox]] reaction, Zn is oxidized to {{chem2|Zn(2+)}} and {{chem2|Cu(2+)}} is reduced to Cu. When electrons are transferred directly from Zn to {{chem2|Cu(2+)}}, the [[enthalpy]] of reaction is lost to the surroundings as heat. However, the same reaction can be carried out in a galvanic cell, allowing some of the chemical energy released to be converted into electrical energy. In its simplest form, a [[half-cell]] consists of a solid metal (called an [[electrode]]) that is submerged in a solution; the solution contains [[cation]]s (+) of the electrode metal and [[anion]]s (−) to balance the charge of the cations.<ref name=":1">{{Cite book|last1=Petrucci|first1=Ralph H.|last2=Herring|first2=F. Geoffrey|last3=Madura|first3=Jeffry D.|last4=Bissonnette|first4=Carey|title=General chemistry: principles and modern applications|publisher=Pearson|year=2017|isbn=978-0-13-293128-1|edition=11th|location=Toronto|pages=869|oclc=951078429}}</ref> The full cell consists of two half-cells, usually connected by a [[semi-permeable membrane]] or by a [[salt bridge]] that prevents the ions of the more noble metal from plating out at the other electrode.<ref name=":1"/>

A specific example is the [[Daniell cell]] (see figure), with  a [[zinc]] (Zn) half-cell containing a solution of {{chem2|ZnSO4}} (zinc sulfate) and a [[copper]] (Cu) half-cell containing a solution of {{chem2|CuSO4}} (copper sulfate). A salt bridge is used here to complete the electric circuit.

If an external electrical conductor connects the copper and zinc electrodes, zinc from the zinc electrode dissolves into the solution as {{chem2|Zn(2+)}} ions (oxidation), releasing electrons that enter the external conductor. To compensate for the increased zinc ion concentration, via the salt bridge zinc ions (cations) leave and sulfate ions (anions) enter the zinc half-cell. In the copper half-cell, the copper ions plate onto the copper electrode (reduction), taking up electrons that leave the external conductor.  Since the {{chem2|Cu(2+)}} ions (cations) plate onto the copper electrode, the latter is called the [[cathode]]. Correspondingly the zinc electrode is the [[anode]]. The electrochemical reaction is:
:{{chem2|Zn(s) + Cu(2+)(aq) -> Zn(2+)(aq) + Cu(s)}}

This is the same reaction as given in the previous example. In addition, electrons flow through the external conductor, which is the primary application of the galvanic cell.

As discussed under [[#Cell voltage|cell voltage]], the [[electromotive force]] of the cell is the difference of the half-cell potentials, a measure of the relative ease of dissolution of the two electrodes into the electrolyte.  The emf depends on both the electrodes and on the electrolyte, an indication that the emf is chemical in nature.

===Half reactions and conventions===
A half-cell contains a metal in two [[oxidation state]]s. Inside an isolated half-cell, there is an [[redox|oxidation-reduction]] (redox) reaction that is in [[chemical equilibrium]], a condition written symbolically as follows (here, "M" represents a metal cation, an atom that has a charge imbalance due to the loss of {{mvar|n}} electrons):
:M{{su|p={{mvar|n}}+|b=(oxidized species)}} + {{mvar|n}} e<sup>−</sup> {{eqm}} M{{sub|(reduced species)}}

A galvanic cell consists of two half-cells, such that the electrode of one half-cell is composed of metal&nbsp;A, and the electrode of the other half-cell is composed of metal&nbsp;B; the redox reactions for the two separate half-cells are thus:
:A<sup> {{mvar|n}}+</sup> + {{mvar|n}} e<sup>−</sup>  {{eqm}}  A
:B<sup> {{mvar|m}}+</sup> + {{mvar|m}} e<sup>−</sup>  {{eqm}}  B

The overall balanced reaction is:
:{{mvar|m}} A + {{mvar|n}} B<sup> {{mvar|m}}+</sup> {{eqm}} {{mvar|n}} B + {{mvar|m}} A<sup> {{mvar|n}}+</sup>

In other words, the metal atoms of one half-cell are oxidized while the metal cations of the other half-cell are reduced. By separating the metals in two half-cells, their reaction can be controlled in a way that forces transfer of electrons through the external circuit where they can do useful [[Work (physics)|work]]. The electrodes are connected with a metal wire in order [[Electrical conduction|to&nbsp;conduct]] the electrons that participate in the reaction.

In one half-cell, dissolved metal&nbsp;B cations combine with the free electrons that are available at the interface between the solution and the metal&nbsp;B electrode; these cations are thereby neutralized, causing them to [[Precipitation (chemistry)|precipitate]] from solution as deposits on the metal&nbsp;B electrode, a process known as [[plating]].

This reduction reaction causes the free electrons throughout the metal&nbsp;B electrode, the wire, and the metal&nbsp;A electrode to be pulled into the metal&nbsp;B electrode. Consequently, electrons are wrestled away from some of the atoms of the metal&nbsp;A electrode, as though the metal&nbsp;B cations were reacting directly with them; those metal&nbsp;A atoms become cations that dissolve into the surrounding solution.

As this reaction continues, the half-cell with the metal&nbsp;A electrode develops a positively charged solution (because the metal&nbsp;A cations dissolve into it), while the other half-cell develops a negatively charged solution (because the metal&nbsp;B cations precipitate out of it, leaving behind the anions); unabated, this imbalance in charge would stop the reaction. The solutions of the half-cells are connected by a salt bridge or a porous plate that allows ions to pass from one solution to the other, which balances the charges of the solutions and allows the reaction to continue.

By definition:
*The [[anode]] is the electrode where oxidation (loss of electrons) takes place (metal&nbsp;A electrode); in a galvanic cell, it is the negative electrode, because when oxidation occurs, electrons are left behind on the electrode.<ref name=r1>{{cite web |title=An introduction to redox equilibria |website=Chemguide (chemguide.co.uk) |url=http://www.chemguide.co.uk/physical/redoxeqia/introduction.html#top |access-date=20 July 2014}}</ref> These electrons then flow through the external circuit to the cathode (positive electrode) (while in electrolysis, an electric current drives electron flow in the opposite direction and the anode is the positive electrode).
*The [[cathode]] is the electrode where reduction (gain of electrons) takes place (metal&nbsp;B electrode); in a galvanic cell, it is the positive electrode, as ions get reduced by taking up electrons from the electrode and plate out (while in electrolysis, the cathode is the negative terminal and attracts positive ions from the solution). In both cases, the statement 'the <u>cat</u>hode attracts <u>cat</u>ions' is true.

By their nature, galvanic cells produce [[direct current]]. 

The ''[[Weston cell]]'' has an anode composed of [[cadmium]] [[mercury (element)|mercury]] [[amalgam (chemistry)|amalgam]], and a cathode composed of pure mercury. The electrolyte is a (saturated) solution of [[cadmium sulfate]]. The [[depolarizer]] is a paste of mercurous sulfate. When the electrolyte solution is saturated, the voltage of the cell is very reproducible; hence, in 1911, it was adopted as an international standard for voltage.

In the strictest sense, a ''battery'' is a set of two or more galvanic cells that are connected [[series circuits|in&nbsp;series]] to form a single source of voltage. For instance, a typical 12&nbsp;V [[lead–acid battery]] has six galvanic cells connected in series, with the anodes composed of lead and cathodes composed of lead dioxide, both immersed in [[sulfuric acid]].

Large central office [[battery room]]s – in a [[telephone exchange]] to provide power for subscribers' [[land-line]] telephones, for instance – may have many cells, connected both in series and parallel: Individual cells are connected in  series as a ''battery'' of cells with some standard voltage ({{circa|40&nbsp;V}}), and ''banks'' of such serial batteries, themselves connected in parallel, to provide adequate amperage to supply a typical peak demand for telephone connections.

===Cell voltage===
The voltage ([[electromotive force]] {{mvar|E}}{{sup|o}}) produced by a galvanic cell can be estimated from the standard [[Gibbs free energy]] change in the electrochemical reaction according to:

<math display="block">\ E_\mathsf{cell}^\mathsf{\ \!o} ~~=\ -\frac{\ \Delta_r G^\mathsf{\ \!o}\ }{\ \nu_\mathsf{e} F\ }\ </math>

where {{mvar|&nu;}}{{sub|e}} is the number of electrons transferred in the balanced half reactions, and {{mvar|F}} is [[Faraday's constant]]. However, it can be determined more conveniently by the use of a [[Table of standard electrode potentials|standard potential table]] for the two [[half cell]]s involved. The first step is to identify the two metals and their ions reacting in the cell. Then one looks up the [[standard electrode potential]], {{mvar|E}}<sup>o</sup>, in [[volt]]s, for each of the two [[half-reaction|half reactions]]. The standard potential of the cell is equal to the more positive {{mvar|E}}<sup>o</sup> value minus the more negative {{mvar|E}}<sup>o</sup> value.

For example, in the figure above the solutions are {{chem2|CuSO4}} and {{chem2|ZnSO4}}. Each solution has a corresponding metal strip in it, and a salt bridge or porous disk connecting the two solutions and allowing {{chem2|SO4(2−)}} ions to flow freely between the copper and zinc solutions. To  calculate the standard potential one looks up copper and zinc's half reactions and finds:

:{{chem2|Cu(2+) + 2e- <-> Cu}} ; {{mvar|E}}<sup>o</sup> = +0.34&nbsp;V
:{{chem2|Zn(2+) + 2e- <-> Zn}} ; {{mvar|E}}<sup>o</sup> = −0.76&nbsp;V

Thus the overall reaction is:
:{{chem2|Cu(2+) + Zn <-> Cu + Zn(2+)}}

The standard potential for the reaction is then {{nobr|+0.34 V − (−0.76 V) {{math|{{=}}}} +1.10 V .}} The polarity of the cell is determined as follows. Zinc metal is more strongly reducing than copper metal because the standard (reduction) potential for zinc is more negative than that of copper. Thus, zinc metal will lose electrons to copper ions and develop a positive electrical charge. The [[equilibrium constant]], {{mvar|K}}, for the cell is given by:

<math display="block">\ \log_e K ~~=~~ \frac{~~~ \nu_\mathsf{e}\;\!F\ E_\mathsf{cell}^\mathsf{\ \!o}\ }{\ R\ T\ }\ </math>

where
: {{mvar|F}} is the [[Faraday constant]],
: {{mvar|R}} is the [[gas constant]], and
: {{mvar|T}} is the [[absolute temperature]] in [[Kelvin (unit)|Kelvins]].
For the Daniell cell {{nobr|{{math| ''K'' ≈ {{val|1.5e37}} }}.}} Thus, at equilibrium, a few electrons are transferred, enough to cause the electrodes to be charged.<ref name=Atkins-dePala-2006>
{{cite book
 |last1=Atkins   |first1=P.
 |last2=de Paula |first2=J.
 |year=2006
 |title=Physical Chemistry |edition=8th
 |publisher=Oxford University Press
 |isbn=978-0-19-870072-2
 |at=Ch.&nbsp;7 and Ch.&nbsp;25, §&nbsp;25.12
}}
</ref>{{rp|style=ama|at=ch.&nbsp;7, "Equilibrium electrochemistry" §§ }}

Actual half-cell potentials must be calculated by using the [[Nernst equation]] as the solutes are unlikely to be in their standard states:

<math display="block">\ E_\mathsf{\;\! half-cell} ~~=~~ E^\mathsf{\ \!o}\ -\ \frac{\ R\ T\ }{\ \nu_\mathsf{e}\ F\ }\ \log_e Q\ </math>

where  {{mvar|Q}}  is the [[reaction quotient]]. When the charges of the ions in the reaction are equal, this simplifies to:

<math display="block">\ E_{\;\! \mathsf{half-cell}} ~~=~~ E^\mathsf{\ \!o}\ -\ 2.303\ \frac{\ R\ T\ }{\ \nu_\mathsf{e} F\ }\ \log_{10} \left\{~~ \mathsf{M}^{n+} \right\}\ </math>

where  M{{sup| {{mvar|n}}+}}  is the [[Activity (chemistry)|activity]] of the metal ion in solution. In practice concentration in {{small|{{sfrac| mol | L }}}} is used in place of activity. The metal electrode is in its standard state so by definition has unit activity. The potential of the whole cell is obtained as the difference between the potentials for the two half-cells, so it depends on the concentrations of both dissolved metal ions. If the concentrations are the same the Nernst equation is not needed, and <math>~ E_\mathsf{cell} ~=~ E_\mathsf{cell}^\mathsf{\;\!o} ~~</math> under the conditions assumed here.

The value of {{nobr|{{math| 2.303 {{sfrac| ''R'' | ''F'' }} }}}} is {{nobr|{{math|1.9845{{small|×}}{{10^|-4}} }} {{small|{{sfrac| V | K }}}} ,}} so at {{math|  ''T'' {{=}} }} 25&nbsp;°C (298.15&nbsp;K) the half-cell potential will change by only {{sfrac|0.05918&nbsp;V | {{mvar|&nu;}}{{sub|e}} }} if the concentration of a metal ion is increased or decreased by a {{nobr|factor of {{math|10}} .}}

<math display="block">\ E_{\;\! \mathsf{half-cell}} ~~=~~ E^\mathsf{\ \!o}\ -\ \frac{\ 0.05918\ \mathsf{V}\ }{\ \nu_\mathsf{e}\ } \log_{10} \left\{~~ \mathsf{M}^{n+} \right\}\ </math>

These calculations are based on the assumption that all chemical reactions are in equilibrium. When a current flows in the circuit, equilibrium conditions are not achieved and the cell voltage will usually be reduced by various mechanisms, such as the development of [[overpotential]]s.<ref name=Atkins-dePala-2006/>{{rp|style=ama|at=§&nbsp;25.12 "Working galvanic cells"}} Also, since chemical reactions occur when the cell is producing power, the electrolyte concentrations change and the cell voltage is reduced. A consequence of the temperature dependency of standard potentials is that the voltage produced by a galvanic cell is also temperature dependent.

==Galvanic corrosion==
{{Main|Galvanic corrosion}}

Galvanic corrosion is the [[electrochemical]] erosion of metals. [[Corrosion]] occurs when two dissimilar metals are in contact with each other in the presence of an [[electrolyte]], such as salt water. This forms a galvanic cell, with hydrogen gas forming on the more noble (less active) metal. The resulting electrochemical potential then develops an electric current that electrolytically dissolves the less noble material. A [[concentration cell]] can be formed if the same metal is exposed to two different concentrations of electrolyte.

==Types==
{{div col begin|colwidth=10em}}
* [[Concentration cell]]
* [[Electrolytic cell]]
* [[Electrochemical cell]]
* [[Lemon battery]]
* [[Thermogalvanic cell]]
{{div col end}}

==See also==
{{div col begin|colwidth=15em}}
* [[Bioelectrochemical reactor]]
* [[Resting potential]]
* [[Bio-nano generator]]
* [[Cell notation]]
* [[Desulfation]]
* [[Electrochemical engineering]]
* [[Electrode potential]]
* [[Electrohydrogenesis]]
* [[Electrosynthesis]]
* [[Enzymatic biofuel cell]]
* [[Galvanic series]]
* [[Isotope electrochemistry]]
* [[List of battery types]]
* [[Sacrificial anode]]
{{div col end}}

==References==
{{reflist|25em}}

==External links==
* [https://web.archive.org/web/20141124233113/http://www.miniscience.com/link/Airbattery.htm How to build a galvanic cell battery] from MiniScience.com
* [http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/galvan5.swf Galvanic Cell], an animation
* [https://web.archive.org/web/20110719210759/http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/electroChem/voltaicCell20.html Interactive animation of Galvanic Cell]. Chemical Education Research Group, Iowa State University.
* [https://av.tib.eu/media/34660?pag=1 Electron transfer reactions and redox potentials in GALVANIc cells - what happens to the ions at the phase boundary (NERNST, FARADAY)] (Video by SciFox on TIB AV-Portal)

{{Galvanic cells}}
{{Authority control}}

[[Category:Galvanic cells| ]]
[[Category:Electrochemical concepts]]
[[Category:Corrosion]]